Editing Chlorine (section)

===Other binary chlorides===
[[File:Nickel(II)-chloride-hexahydrate-sample.jpg|thumb|right|Hydrated [[nickel(II) chloride]], NiCl<sub>2</sub>(H<sub>2</sub>O)<sub>6</sub>]]
Nearly all elements in the periodic table form binary chlorides. The exceptions are decidedly in the minority and stem in each case from one of three causes: extreme inertness and reluctance to participate in chemical reactions (the [[noble gas]]es, with the exception of [[xenon]] in the highly unstable [[xenon dichloride|XeCl<sub>2</sub>]] and XeCl<sub>4</sub>); extreme nuclear instability hampering chemical investigation before decay and transmutation (many of the heaviest elements beyond [[bismuth]]); and having an electronegativity higher than chlorine's ([[oxygen]] and [[fluorine]]) so that the resultant binary compounds are formally not chlorides but rather oxides or fluorides of chlorine.<ref name="Greenwood821">{{harvnb|Greenwood|Earnshaw|1997|pp=821–844}}</ref> Even though [[nitrogen]] in NCl<sub>3</sub> is bearing a negative charge, the compound is usually called [[nitrogen trichloride]].

Chlorination of metals with Cl<sub>2</sub> usually leads to a higher oxidation state than bromination with Br<sub>2</sub> when multiple oxidation states are available, such as in [[molybdenum(V) chloride|MoCl<sub>5</sub>]] and [[molybdenum(III) bromide|MoBr<sub>3</sub>]]. Chlorides can be made by reaction of an element or its oxide, hydroxide, or carbonate with hydrochloric acid, and then dehydrated by mildly high temperatures combined with either low pressure or anhydrous hydrogen chloride gas. These methods work best when the chloride product is stable to hydrolysis; otherwise, the possibilities include high-temperature oxidative chlorination of the element with chlorine or hydrogen chloride, high-temperature chlorination of a metal oxide or other halide by chlorine, a volatile metal chloride, [[carbon tetrachloride]], or an organic chloride. For instance, [[zirconium dioxide]] reacts with chlorine at standard conditions to produce [[zirconium tetrachloride]], and [[uranium trioxide]] reacts with [[hexachloropropene]] when heated under [[reflux]] to give [[uranium tetrachloride]]. The second example also involves a reduction in [[oxidation state]], which can also be achieved by reducing a higher chloride using hydrogen or a metal as a reducing agent. This may also be achieved by thermal decomposition or disproportionation as follows:<ref name="Greenwood821" />
: EuCl<sub>3</sub> + {{sfrac|1|2}} H<sub>2</sub> ⟶ EuCl<sub>2</sub> + HCl
: ReCl<sub>5</sub> {{overunderset|⟶|at "bp"|&thinsp;}} ReCl<sub>3</sub> + Cl<sub>2</sub>
: AuCl<sub>3</sub> {{overunderset|⟶|160&thinsp;°C|&thinsp;}} AuCl + Cl<sub>2</sub>

Most metal chlorides with the metal in low oxidation states (+1 to +3) are ionic. Nonmetals tend to form covalent molecular chlorides, as do metals in high oxidation states from +3 and above. Both ionic and covalent chlorides are known for metals in oxidation state +3 (e.g. [[scandium chloride]] is mostly ionic, but [[aluminium chloride]] is not). [[Silver chloride]] is very insoluble in water and is thus often used as a qualitative test for chlorine.<ref name="Greenwood821" />